How electron configuration explains trends across the periodic table
The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. The long form of the table has 18 groups and 7 periods and is divided into s, p, d and f blocks according to which subshell receives the last electron. Elements in the same group share the same outer electron configuration and similar chemistry. Elements with atomic number above 100 are given temporary IUPAC names based on their digits.
Several properties follow clear trends. Atomic radius decreases across a period as nuclear charge increases, and increases down a group as new shells are added. Ionisation enthalpy generally increases across a period, but there are exceptions; nitrogen has a higher value than oxygen because of its stable half-filled p subshell. Electron gain enthalpy becomes more negative across a period, and chlorine has a more negative value than fluorine because fluorine is small and crowded. Electronegativity increases across a period and decreases down a group. Metallic character decreases across a period. Diagonal relationships link pairs like lithium and magnesium.
Why it matters
Periodic trends let chemists predict the reactivity, bonding and properties of elements without testing each one, and they helped scientists predict elements before they were discovered. Trends explain why alkali metals are stored in oil and why fluorine is the most reactive non-metal. This topic is used throughout bonding, the p-block and coordination chemistry.
Key terms
- Ionisation enthalpy
- The energy needed to remove the most loosely held electron from a gaseous atom.
- Electron gain enthalpy
- The energy change when an electron is added to a neutral gaseous atom.
- Electronegativity
- The tendency of an atom in a bond to attract shared electrons towards itself.
- Diagonal relationship
- The similarity between certain elements diagonally placed in the second and third periods.